Lawrence Berkeley National Laboratory Recent Work

Lawrence Berkeley National Laboratory Recent Work Title GEOCHEMICAL AND COSMOCHEMICAL CYCLES INVOLVING SULFUR, SULFIDE, SULFITE, AND SULFATE Permalink https://escholarship.org/uc/item/3jb1c1qd Author Meyer, B. Publication Date 1979-07-01 eScholarship.org Powered by the California Digital Library University of California LBL-9396 (!_~ Preprint · lrnl Lawrence Berkeley Laboratory ~ UNIVERSITY OF CALIFORNIA, BERKELEY, CA Materials & Molecular Research Division Submitted to Geochimica Cosmochimica Acta GEOCHEMICAL AND COSMOCHEMICAL CYCLES INVOLVING SULFUR, SULFIDE, SULFITE, AND SULFATE -. ECEP/ED B. Meyer, L. Peter and M. Ospina ~o./l.WR!:NCE Sl:fi..KW,'ZY I.A30RAiORY ! IJI 7 Q HY"Ig July 1979 .J t L u 1::1/ LlBR!.RY AND DOCUMENTS SE:Cl'IO.N TWp-WEEK LOAN COPY This is a Library Circulating Copy which may be borrowed for two weeks. For a personal retention copy, call Tech. Info. Division, Ext. 6782 Prepared for the U. S. Department of Energy under Contract W-7405-ENG-48 DISCLAIMER This document was prepared as an account of work sponsored by the United States Government. While this document is believed to contain correct information, neither the United States Government nor any agency thereof, nor the Regents of the University of California, nor any of their employees, makes any warranty, express or implied, or assumes any legal responsibility for the accuracy, completeness, or usefulness of any information, apparatus, product, or process disclosed, or represents that its use would not infringe privately owned rights. Reference herein to any specific commercial product, process, or service by its trade name, trademark, manufacturer, or otherwise, does not necessarily constitute or imply its endorsement, recommendation, or favoring by the United States Government or any agency thereof, or the Regents of the University of California. The views and opinions of authors expressed herein do not necessarily state or reflect those of the United States Government or any agency thereof or the Regents of the University of California. 0 GEOCHEMICAL AND COSMOCHEMICAL CYCLES INVOLVING SULFUR, SULFIDE, SULFITE AND SULFATE B. Meyer, * L. Peter and M. Ospina Department of Chemistry University of Washington Seattle, Washington 98195 and Molecular and Materials Research Division Lawrence Berkeley Laboratory University of California Berkeley, California 94720 ABSTRACT Raman spectra· of aqueous systems containing sulfur dioxide, elemental sulfur and sulfate indicate that the equilibrium between these species is catalyzed by elemental sulfur. There fore, dynamic equilibrium can be expected under conditions prevalent on Venus, on Io and in epigenic sulfur deposits. Sulfur, sulfuric acid, sulfur dioxide and water are in contact with each other in a variety of systems such as planetary atmospheres (Knollenberg, 1979; Nelson, 1978; Oyama, 1979; and Stewart, 1979), bio-epigenic sulfur deposits (Rucknick, 1979) as well as in industrial so2 recovery systems in coal burning power plants (Meyer, 1977). The reactions and equilibria between elemental sulfur, sulfur dioxide, sulfuric acid and water can be summarized in two over-all equations: Eq. 3S02 + 2H 20 2H 2so4 + 1/n sn + Eq. I I 1/n sn + 4H 2o 3H 2S H2so4 At room temperature pure liquid so2 in contact with water is stable for many years. Likewise, elemental sulfur, water and sulfuric acid do not readily react. However, contrary to· the widely held belief, this sulfur system is not inert. Priestley, 1970, heated sulfur dioxide in a sealed tube and observed the reaction over a period of several months and observed crystals which he identifi.ed as elemental sulfur. Boehm, 1882, studied reaction Eq. II. Bichowsky, 1922, studied both systems carefully and discovered that the reactions are rever sible. Despite its importance and practical applications, the chemistry of this system has not been studied in more than half a century, even though Thode's pioneering work on isotope fractionization (Monster, 1965) stimulated interest among earth scientists (Ellis, 1971; Dana, 1966; Dana 1971; and Robertson, 1973). Unfortunately, thermodynamics are only of limited help, because of the uncertainty about the variety of states in which the various sulfur species can appear. We report here some qualitative data based on experiments in which we used high purity elemental sulfur and ~situ observation techniques. Neither was available to earlier authors. At room temperature, we confirm that the pure chemicals in reaction (I) do not react notice ably over a period of at least five years. However, in the presence of ammonia or other suitable catalysts, reaction commences within weeks (Meyer, 1979), elemental sulfur precipi tates from tubes containing sulfur dioxide within one to three years, and the resulting -1- 0 r H20 -' 5042 b_;:~ HS04 s H0 2 2 c so.;- - ~ d ~ ~20 ~ XBL 793-8912 Fig. 1. Raman spectra of sulfur systems: (a) S02 and H20 after 27 hours at 80°C, (b) elemental sulfur and water after 24 hours at 190°C, (c) elemental sulfur, S02 and H20 after 12 hours at 120°C, and (d) so2 and water with 0.1 MBac1 2 after 12 hours at 25°C. solution contains dithionate and other sulfur oxyacids as transient species, in addition to sulfate which constitutes the most stable end product. Figure la shows the Ramen spectrum of the aqueous phase in a sealed pyrex tube 6 mm diam. x 6 em containing 0.2 ml. H20 and 0.2 ml liquid sulfur dioxide at 25°C after it was heated for 27 hrs at 80°C. The liquid remained clear and colorless. The broad peaks at 470 cm-1 and 800 cm-1 are characteristic for colloidal sulfur; the weak shoulders at 980 cm-l and 1,050 cm-l are due to sulfate and bisul 1 fate, respectively. The strongest peak at 1,144 cm- , due to unreacted sulfur dioxide, is ) not shown. The Raman band intensities can be evaluated by comparison with the internal water ·standard (Meyer, 1979) and correspond to 0.25 M sulfate and 0.30 M bisulfate solutions. Our observations are consistent with Bichowsky's equilibrium studies (Bichowsky, 1922), however, we have found that at 80°C reaction (I) proceeds about halfway in one day or less, while Bichowsky observed that at 312°C equilibrium in the opposite direction required at least three to seven days. -2- According to earlier work (Bichowsky, 1922) reaction (II) takes sixty days to reach equili brium at 260°C. However, using high purity sulfur, we observed already after 24 hrs at 190°C the Raman spectrum in Figure lb. The main peaks correspond to 0.01 M sulfate (so -2) and 4 0.03 M bisulfide (HS-1). Upon further heating, the concentration of sulfuric acid does not increase rapidly, but hydrogen sulfide disappears, probably because sulfuric acid and elemen tal sulfur enter reaction (I), and the resulting sulfur dioxide consume hydrogen sulfide yielding elemental sulfur. The latter reaction is well known and was first described by Wackenroder, 1846. It involves the formation of intermediate species in intermediate oxida tion states, including thiosulfate, tetrathionate, and other polythionates which are not stable under our conditions (Meyer, 1977). The existence of hydrogen sulfide in reaction (II) suggests that hydrogen sulfide might also occur as an intermediate in reaction (I). Unfortunately, this cannot be easily verified, beca·use of the above mentioned Wachenroder reaction. In further experiments, we added elemental sulfur to the reagents in reaction (I). Figure lc shows the Raman spectrum of this system after it was kept for 12 hrs at 120°C. The peaks at 1 1 1 1,050 cm- and 980 cm- correspond to 1 M Hso4- solution and 0.2 M sulfate. Obviously, this reaction is far faster than with so2 and water alone, as observed in the experiment shown in Figure la. We observed the behavior of this system over a range of temperatures between 20°C and 200°C, and conclude that it is catalyzed by solid as well as liquid elemental sulfur. We do not understand the mechanism, but the affinity between sulfur dioxide and elemental sulfur is consistent with our earlier observation that the solubility of elemental sulfur in so2 exhibits unusual temperature dependence (Meyer, 1971). The solubility and catalysis effects are both difficult to measure quantitatively, because sulfur undergoes a phase tran sition at 95°C and changes its allotropic composition above the melting point at 124°C. The catalysis appears to be inhibited by a factor of four in 1M HCl, but this effect is due to the fact that HCl reduces the solubility of S02 in water. In the last set of experiments we sealed so2 in a tube containing 0.1 M BaC1 2. The solution promptly clouded at room temperature due to the formation of BaS03 which slowly precipitated. Figure ld shows the Raman spectrum of the suspension after 9 hrs. The strongest peak at 960 cm-1 is due to barium sulfite. The peaks at ?00, 490 and 460 cm-1 are due to orthorhom bic elemental sulfur. The peak at 980 cm-1 corresponds to sulfate. Since barium sulfate has a solubility product (Cohen, 1940) of about lo-10 , our observation indicates that the sulfate ion concentration reaches at least 10-9 M in our system. Our experiments show that the ubiquitous reactions (I) and (II) can proceed over a wide range of temperatures, and thus can take a dynamic role wherever elemental sulfur and and sulfuric acid or sulfite and water are present. The Pioneer experiments (Knollenberg, 1979; Oyama, 1979; Stewart, 1979) have recently indicated that in the upper cloud layer in the atmosphere of Venus at 13°C, sulfuric acid (Knollenberg, 1979) and sulfur dust are present. In lower, hotter levels, 30 miles above the ground, sulfur dioxide (Oyama, 1979; Stewart, 1979) and water vapor at about 200°C are present.

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